introduction-to-inorganic-chemistry

Introduction to Inorganic Chemistry	Loose compounds and solutions

Note

A “normal” solution used to refer to its concentration, but this usage has now been dropped.

14.1.2 Abnormal solutions

These are where solute, solvent, and solution do not all have the same character. Examples are:

•solutions of salt-like substances in molecular ones (e.g. sodium chloride in water);

•solutions of molecular substances having similar properties to the previous type (e.g. hydrogen chloride in water);

•solutions of metals in molecular substances (e.g. sodium in liquid ammonia).

I will consider these in turn.

Solutions of salt-like substances in molecular ones

Sodium chloride dissolves in water despite the fact that the latter is molecular. The solutions are good conductors of electricity, and undergo electrolysis in the same way as molten sodium chloride except that hydrogen is formed at the cathode (sodium reacts with water to give hydrogen). The freezing points of the solutions are approximately twice those expected for the presence of NaCl molecules (the factor approaches 2.00 at low concentrations). The solutions evidently contain Na+ and Cl− ions like sodium chloride itself (Chap. 5).

Given the strong Coulombic attraction between Na+ and Cl− ions in solid sodium chloride, it is at first sight surprising that the latter should dissolve in water. The reason must be that the H2O molecule is very polar. We saw in Chapter 5 that the charge distribution is approximately (H0.4+)2O0.8−. Thus if H2O molecules surround the Na+ ions with O0.8− pointing towards them and the Cl− ions with H0.4+ pointing towards them, considerable Coulombic attraction is generated, enough presumably to dissolve the ions.

For some cations, this attraction may be supplemented by some degree of dative bonding. For example, violet solutions of chromium(III) salts behave as if they contain relatively tightly bound [Cr(H2O)6]3+ ions, analogous to [Cr(NH3)6]3+ ions. Thus, if a solution containing a Cr3+ ions and b H2O molecules is added to water enriched with H218O and some water is distilled from the mixture, the 18O:16O ratio in this corresponds to rapid exchange, not with all the H2O molecules in the original solution, but with (b − 6a).

Solutions of salt-like substances do not always contain simple hydrated ions. For example, some solutions of chromium(III) salts are green. These contain complexes with the anions of the salt, e.g. [CrCl2(H2O)4]+. Freezing points of copper sulfate solutions indicate that, except at very low concentrations, most of the salt is present, not as Cu2+ and SO42− ions, but as [CuSO4(H2O)n]:

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Introduction to Inorganic Chemistry Loose compounds and solutions

Concentration	Proportion as [CuSO4aq]
0.001 M	15%
0.01 M	50%
0.1 M	80%

When solutions containing only simple hydrated ions are required, salts of weakly coordinating cations

(e.g. K+) or anions (e.g. ClO4−, BF4−) must be used.

Water is a particularly good ionizing solvent, but there are others. Among the most important are liquid NH3 (b.p. −33 °C), liquid HF (b.p. 19.5 °C), 100% CH3CO2H (b.p. 118 °C), 100% H2SO4 (b.p. 320 °C), and 100% HSO3F (b.p. 163 °C).

Salt-like solutions of molecular substances

Anhydrous hydrogen chloride is a gas. However, it dissolves in water to give solutions that are very like sodium chloride solutions. That is, they conduct electricity and undergo electrolysis, giving hydrogen at the cathode and chlorine at the anode. Their freezing points are also approximately twice those expected for the presence of HCl molecules. The solutions evidently contain hydrated H+ and Cl− ions. They will be discussed further in Chapter 17.

Solutions of metals in liquid ammonia

Sodium dissolves in liquid ammonia. Dilute solutions are blue, concentrated ones are bronze-coloured. The solutions conduct electricity. The other alkali metals and the alkaline earth metals also dissolve in liquid ammonia, giving solutions of the same colour.

Since sodium metal comprises Na+ ions and electrons (Chap. 5), the solutions may contain solvated Na+ ions and electrons. This is supported by the fact that other metals give solutions of the same colour, the colour being due to solvated electrons. The formation of the latter seems to be due to the slowness of the expected reaction:

Na + NH3 → NaNH2 + ½H2

The solutions are indeed metastable, and catalysts soon give NH2− + ½H2.

14.2 Loose compounds

These are formed between the components of some solutions. They may be divided into two classes: (i) molecule-molecule compounds; (ii) salt-molecule compounds.

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Introduction to Inorganic Chemistry	Loose compounds and solutions

14.2.1 Molecule-molecule compounds

When chlorine is passed into water cooled in ice, pale greenish-yellow crystals of chlorine hydrate separate. Chemical analysis gives its composition as approximately Cl2·7H2O. An X-ray examination of its structure gives its ideal formula as 3Cl2·23H2O or Cl2·7.67H2O. When gently warmed the crystals melt and in the dark evolve pure chlorine. The association between the chlorine molecules and the water molecules is evidently very weak.

Similar compounds are formed with other solutes and solvents, and between other molecular compounds generally.

Bonding

Compounds of this type seem generally to be held together by means of the same sort of forces by which the parent compounds are held together, viz. van der Waals’ forces, plus forces arising from any polarity that the molecules may have. In some cases, however, these forces may be supplemented by varying degrees of dative bonding, leading in the limit to compounds of the coordination type.

As we have seen, the water molecule has considerable polarity, whence the low volatility of water. In a compound like 3Cl2·23H2O, the main force operating seems to be the attraction between the dipoles of the water molecules, forming a cage around the chlorine molecules. Such compounds are called “clathrate” compounds (Latin clathratus, meaning “enclosed by cross bars of a grating”).

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Introduction to Inorganic Chemistry	Loose compounds and solutions

14.2.2 Salt-molecule compounds

The best-known examples of these are some of the salt hydrates. For example, sodium sulfate crystallizes from water at room temperature in the form of large crystals of Na2SO4·10H2O. These readily lose their water of crystallization, efflorescing in air to a powder of anhydrous Na2SO4. The water molecules are evidently only relatively loosely associated with the ions of the salt.

Not all salt hydrates are, of course, loose compounds. For many of them the association of at least some of the water molecules is much stronger. An instructive example is provided by the compound CrCl3·6H2O, which exists in three forms.

(i)A dark green form, made by crystallizing the anhydrous chloride from water. One mole of this loses two moles of water when placed over concentrated sulfuric acid, and precipitates only one mole of AgCl to start with when treated with silver nitrate solution.

(ii)A greyish-blue form, made by refluxing a solution of (i) for some time, then cooling the solution in ice and saturating it with hydrogen chloride. This does not lose water over sulfuric acid, and immediately precipitates all of its chloride on treatment with silver nitrate.

(iii)A light green form, made by adding ether saturated with hydrogen chloride to the mother liquor from the preparation of (ii) and passing in hydrogen chloride. One mole of this loses one mole of water over sulfuric acid, and immediately precipitates two moles of AgCl with silver nitrate.

The properties of (i) - (iii) suggest the following formulations:

(i)[CrCl2(H2O)4]Cl·2H2O

(ii)[Cr(H2O)6]Cl3

(iii)[CrCl(H2O)5]Cl2·H2O

All three contain coordinated water molecules; (i) and (iii) contain loosely held water as well. CrCl3 also forms a brown, ether-soluble trihydrate, [CrCl3(H2O)3]0.

Similar compounds to salt hydrates are formed with other solvents. In all cases, a broad distinction can be drawn between coordinated solvent molecules and loosely associated ones, but the line between the two categories is not a sharp one.

Bonding

Loose compounds of the salt-molecule type seem to depend for their formation on the polarity of the molecular component, enabling it to be attracted electrostatically to the ions of the salt (positive poles to anions, negative to cations). This attraction can vary in strength, leading to the spectrum of types from loose compounds to coordination compounds indicated above.

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Introduction to Inorganic Chemistry	Loose compounds and solutions

As discussed for salt-molecule solutions, the forces of attraction between the ions and the molecules cannot be trivial since the ions have to move away from each other to make room for the molecules, and this entails a loss of Coulombic energy. The formation of salt hydrates must be due to the high polarity of the water molecule, producing strong forces of attraction between ions and molecules. A “loose” compound is not so much one in which the bonding between components is weak, but one in which the net bonding is weak (i.e. the difference in energy between reactants and products is low).

Nomenclature of loose compounds

Loose compounds are named as illustrated below:

Na2SO4·10H2O

sodium sulfate decahydrate

or sodium sulfate—water (1/10)

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Introduction to Inorganic Chemistry	Types of chemical reaction

15 Types of chemical reaction

15.1 General types

1 Combination

This is when two or more elements or compounds unite to form a single compound. For example: 2Mg + O2 → 2MgO

KCN + S → KSCN

K2O + SO3 → K2SO4

2 Decomposition

This is when a compound breaks up into two or more elements or simpler compounds. For example: 2H2O2 → 2H2O + O2

NH4NO3 —heat→ 2H2O + N2O

2FeSO4 —heat→ Fe2O3 + SO3 + SO2

A reversible decomposition is called a “dissociation”. Examples:

PCl5PCl3 + Cl2

2HIH2 + I2

NH4ClNH3 + HCl

3 Displacement

This is when one element replaces another in a compound, or one compound replaces another in a larger compound. For example:

Fe2O3 + 2Al → 2Fe + Al2O3

Zn + CuSO4 —aq→ Cu + ZnSO4

4HNO3 + P4O10 → 4HPO3 + 2N2O5

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Introduction to Inorganic Chemistry	Types of chemical reaction

4 Double decomposition

This is when two compounds interact by exchange of parts to produce two new compounds. The parts may be atoms or radicals. Examples:

AgNO3 + NaCl —aq→ AgCl + NaNO3

PCl3 + 3AgF → PF3 + 3AgCl

SnCl4 + 4PhMgCl → SnPh4 + 4MgCl2

5 Addition

This is combination viewed from the point of view of one of the reactants, which has further atoms or groups of atoms added to it. For example:

C2H4 + Br2 → C2H4Br2

2PCl3 + O2 → 2POCl3

CuSO4 + 5H2O → CuSO4·5H2O

In each example, the first reactant is said to have atoms or groups of atoms from the second added to it.

6 Substitution

This is displacement or double decomposition viewed from the point of view of one of the reactants, in which one atom or radical is replaced by another atom or radical. For example:

CH4 + Cl2 → CH3Cl + HCl

C6H6 + HNO3 → C6H5NO2 + H2O

These are double decompositions:

CH3–H + Cl–Cl → CH3–Cl + H–Cl

C6H5–H + HO–NO2 → C6H5–NO2 + H–OH

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Introduction to Inorganic Chemistry	Types of chemical reaction

7 Insertion

This is an addition reaction in which an atom or group of atoms is inserted between two atoms initially bound together. An example is the reaction

which takes place when R is one of a number of transition-metal radicals, e.g. Mn(CO)5.

8 Isomerization

This is when a substance changes into another form with different properties but the same molecular formula. For example:

CH2=CH–CH2–CH3 —catalyst→ CH3–CH=CH–CH3

red HgI2 —heat→ yellow HgI2

The different forms are called “isomers” (Greek isos, equal), and the phenomenon of the existence of different forms “isomerism”.

A rapidly reversible isomeric change is called a “tautomeric” change. The different forms are called “tautomers”, and the phenomenon “tautomerism”.

An isomeric change in the case of an element is called an “allotropic” change. The different forms are called “allotropes”, and the phenomenon “allotropy”.

9 Polymerization

This is when a substance changes into another substance with the same composition but a much higher molecular mass. For example:

nC2H4 —heat, pressure, catalyst→ (–CH2–CH2–)n

The product of such a reaction is called a “polymer”, and the starting material the corresponding “monomer” (Greek polus, much; meros, share; monos, alone).

The term polymerization is also used for processes in which a polymer is formed, not from the monomer, but from other reactants of low molecular mass. This usage is somewhat misleading, but is well established.

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Introduction to Inorganic Chemistry	Types of chemical reaction

10 Oligomerization

This is similar to polymerization except that the product contains only a small number of monomer units (Greek oligos, small). For example

3C2H2 —heat, catalyst→ C6H6

The degree of oligomerization is specified by the numerical prefixes di, tri, etc., as in dimer, trimerize, etc.

15.2 Some special types of reaction

There are many special types of reaction. Among the more important are the following.

Precipitation

This is when an insoluble solid is formed in a reaction taking place in solution. For example

NaCl + AgNO3 —aq→ AgCl↓ + NaNO3

K2SO4 + BaCl2 —aq→ BaSO4↓ + 2KCl

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Introduction to Inorganic Chemistry	Types of chemical reaction

These are both double decompositions.

Precipitation reactions are popular in analytical chemistry, providing both characteristic tests for different ions, and also, when the precipitate is sufficiently insoluble, a means of determining the amount (“gravimetric analysis”).

Neutralization

This is a reaction between an acid and a base. It is discussed under “Acids, bases and salts” (Chap. 16).

Hydrolysis

This is a double decomposition involving water as one of the reactants. This splits into H and OH or 2H and O. Examples:

POCl3 +3H2O → 3HCl + PO(OH)3

CH3CO2CH3 + H2O → CH3CO2H + CH3OH

SOCl2 + H2O → 2HCl + SO2

Redox reactions

These are discussed separately under the heading “Oxidation and reduction” (Chap. 17).

Solvation

This is an addition reaction between a solute and a solvent. For example:

Na2SO4 + 10H2O → Na2SO4·10H2O

CuSO4 + 5H2O → CuSO4·5H2O

MgBr2 + 2Et2O → MgBr2·2Et2O

The product of such a reaction is called a “solvate”. If the solvent is water the adduct is called a “hydrate”, and the process “hydration”.

Complexation

This is an addition reaction that leads to the formation of a coordination entity (Chap. 13).

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